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1. Group
Presentation
Guru Sai Bhaskari-2301AI19
Neha Reddy Sabbidi-2301AI42
Preeti Kumari-2301AI17
Shiksha Raginee-2301AI13
Prepared by group 4

2. Primary
Bonding

3. Introduction to
bonding
• Bonding is the force of attraction that holds
atoms together in a compound or material.
• Atoms bond to achieve stability by filling their
outermost electron shells.

4. Importance of bonding
• Determines the physical and chemical properties of
materials.
• Essential in the formation of solids, liquids, and
gases.
• Covalent bonds hold DNA, proteins, and enzymes
together.
• Hydrogen bonding helps in the structure of DNA
(double helix).
• Strong bonds in metals and alloys are used in
construction and manufacturing.
• Covalent bonds in polymers are used to make
plastics and synthetic fibers.

5. Types of bonding
• Ionic Bonding
• Covalent Bonding
• Metallic Bonding
Primary Bonds Secondary
Bonds
• Van der Waals
Forces
• Hydrogen Bonding

6. Primary
Bonding Definition:
Primary bonding refers to the strong
interatomic forces that hold atoms
together in a substance. These bonds
involve the transfer, sharing, or
delocalization of electrons and are
responsible for the stability and
structure of materials

7. Ionic Bonding
Ionic bonding is a type of chemical bond that forms between metal and non-metal
atoms when electrons are transferred from one atom to another. This transfer
creates oppositely charged ions that are held together by electrostatic attraction.
Electron Transfer:
• A metal atom loses electrons to become a positively charged ion (cation).
• A non-metal atom gains those electrons to become a negatively charged ion
(anion).
Formation of Ionic Compounds:
• The oppositely charged ions attract each other due to electrostatic forces,
forming a strong bond.
• The result is a crystalline structure, commonly seen in salt (NaCl).
Definition:
How Does Ionic Bonding Work?

8. • Sodium (Na): A metal with 1 electron in its outer shell. It loses this electron to become Na⁺.
• Chlorine (Cl): A non-metal with 7 electrons in its outer shell. It gains 1 electron to become C
• Na⁺ and Cl⁻ attract each other, forming a stable ionic compound (NaCl - Table Salt).
Example:
Formation of
Sodium
Chloride
(NaCl)

9. Properties of Ionic Compounds:
Common Examples of Ionic Compounds:
• High Melting and Boiling Points – Strong ionic bonds require a lot of
energy to break.
• Hard but Brittle – Ionic crystals are hard but can shatter if force is
applied.
• Soluble in Water – Ionic compounds dissolve in water, breaking into ions.
• Conduct Electricity in Solution or Molten State – Ions are free to move
in liquid form, allowing electrical conduction.
• Sodium Chloride (NaCl) – Table salt
• Calcium Fluoride (CaF₂) – Used in toothpaste
• Magnesium Oxide (MgO) – Used in ceramics

11. Covalent
Bonding
Covalent bonding is a type of chemical bond where two or more non-
metal atoms share electrons to achieve stability. Instead of transferring
electrons (like in ionic bonding), atoms share electrons to complete
their outer shells.
• Atoms want to achieve a full outer electron shell (usually 8
electrons, following the Octet Rule).
• Since non-metals have similar electronegativities, they cannot easily
gain or lose electrons. Instead, they share electrons.
• This sharing creates a strong bond between the atoms.
Definition:
How Does Covalent Bonding Work?

12. • Oxygen (O): Needs 2 more electrons to complete its outer shell.
• Hydrogen (H): Needs 1 more electron to be stable.
• Each hydrogen atom shares 1 electron with oxygen, forming two covalent bonds.
• The result is H₂O (water), a stable molecule.
Example:
Formation of
a Water
Molecule
(H₂O)

13. Properties of Covalent Compounds:
Common Examples of Covalent
Compounds:
• Low Melting and Boiling Points – Weak forces between molecules make them easy to
separate.
• Poor Conductors of Electricity – No free-moving charged particles (except in some special
cases like graphite).
• Exist as Gases, Liquids, or Soft Solids – Unlike ionic compounds, many covalent compounds
exist in gaseous or liquid form.
• Insoluble or Partially Soluble in Water – Some covalent compounds do not dissolve in water
(e.g., oil), while others do (e.g., sugar).
• Oxygen (O₂) – Essential for respiration
• Carbon Dioxide (CO₂) – Used by plants in photosynthesis
• Methane (CH₄) – A major component of natural gas
• Diamond (C) & Graphite (C) – Both made of carbon but with
different structures

14. Examples of Covalent compounds in
Everyday life

15. Metallic
Bonding
Metallic bonding is the strong force of attraction between positively charged
metal ions and a "sea" of free-moving electrons in a metal. Unlike ionic or
covalent bonds, where electrons are either transferred or shared between
specific atoms, metallic bonding allows electrons to move freely throughout the
entire structure.
• Metal atoms release their outermost electrons, forming positive metal ions
(cations).
• These free electrons move randomly throughout the metal structure,
creating a "sea of electrons".
• The electrostatic attraction between the positive ions and the free-moving
electrons holds the metal together.
Definition:
How Does Metallic Bonding Work?

16. • Copper atoms lose their valence electrons, becoming Cu²⁺ ions.
• The free electrons move between the ions, binding them together.
• This allows copper to conduct electricity and be shaped easily.
Example: Metallic
Bonding in Copper (Cu)

17. Properties of Covalent Compounds:
High Electrical and Thermal Conductivity
• The free-moving electrons allow metals to conduct electricity and heat efficiently.
• Example: Copper (Cu) is used in electrical wiring.
Malleability and Ductility
• Metals can be hammered into sheets (malleability) and drawn into wires (ductility) because the atoms can slide
over each other without breaking the bond.
• Example: Aluminum (Al) is used in foil and wiring.
High Melting and Boiling Points
• The strong attraction between the metal ions and free electrons requires a lot of energy to break.
• Example: Iron (Fe) has a high melting point, making it useful in construction.
Luster (Shiny Appearance)
• The free electrons reflect light, giving metals their shiny look.
• Example: Gold (Au) and Silver (Ag) are used in jewelry.
• Strength and Hardness
• Metallic bonds make metals strong and durable.
• Example: Steel (a mixture of iron and carbon) is used in bridges and buildings.

18. Examples of Metallic bonding in
Everyday life

20. Conclusion
• Primary bonds are essential for material properties.
• Understanding them helps in material selection and engineering.
• Applications range from everyday items to advanced technologies

21. Thank
you